# Formal Charge, cont.

Recall that Formal Charge (FC) = the charge an atom would have if all bonding pairs were shared equally (polar bonds don't exist in this model).

#### To assign Formal Charges:

1. Draw a correct Lewis Structure.
2. Assign both electrons of a lone pair to its associated atom.
3. Divide all bonding pairs, giving one electron of each pair to each atom in the bond.
4. Calculate FC = # electrons on the unbonded (elemental) atom - # electrons assigned to the bonded atom.

#### Examples:

• Perchlorate ion (ClO4-)
• LS:
• FCCl: 7 - 4 = +3
• FCO: 6 - 7 = -1
• FC = +3 + 4 (-1) = -1. Notice that the Formal charges add up to give -1, the charge on the molecular ion.
• Note can also determine FC's for compounds containing transition metals. Just don't worry about transition metal, use octet rule for representative elements, and find transition metal FC by difference. For example, permanganate ion MnO4-
• Oxygen will have an octet in normal oxygen compounds, and will share one bonding pair, thus
• FCO: 6 - 7 = -1
• FC = FCMn + 4 (FCO) = -1, the charge on the ion, thus
• FCMn = -1 - 4 (FCO) = -1 - 4 (-1) = +3

## Lewis Structures for Covalent Molecules-Octet Variations

### Multiple Bonds & Resonance

Recall we must show an octet (or duet for Period I) in the outer-most shell (valence electrons). When this does not occur with single electron pairs (bonds) between atoms can sometimes make it happen with multiple bonds. You might find "Clark's Method" useful for determining the bonding patterns of various molecules:

• #### Clark's Method (abbreviated)for determining bonding in covalent Lewis Structures:

• Add up all of the valence electrons in the structure (remember to add one electron for each negative charge, or subtract one for each positive charge)
• If e- = 6y + 2 where y = # atoms other than H, then octet rule is followed with single bonds only.
• If e- < 6y + 2 then probably have multiple bonding with the number of multiple bonds = /2 (remember a triple bond is 2 multiple bonds!).
• If e- > 6y + 2 then have an expanded valence shell.
• If you can draw more than one structure, then chose the most symmetrical.
• If two or more structures are equally symmetrical, then you probably have resonance and should show all structures connected by double arrows.
• Multiple bond examples:
• Carbon monoxide, CO
• valence electrons = 4 + 6 = 10
• 6y + 2 = 14, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
• LS = :C:::O:
• Carbon dioxide, CO2
• valence electrons = 4 + 2x6 = 16
• 6y + 2 = 20, thus 4 fewer electrons than required for all single bonds, 4/2 = 2 multi-bonds (2 double or 1 triple)
• LS: from symmetry C will be central atom, therefore=

Practice example:

• Carbon disulfide
• Hint: sulfur is immediately below oxygen in Periodic Table, so it should be similar to carbon dioxide.
• Resonance example:
• Carbonate ion - CO32-
• valence electrons = 4 + 3 (6) + 2 = 24
• 6y + 2 = 26, but e- = 24, therefore expect one multiple bond.
• LS =
• However, other equally symmetrical structures are possible, so:
• Expanded Valence Shell Example:
• SF4
• valence electrons = 6 + 4(7) = 34
• 6y + 2 = 32, but e- = 34, therefore expect expanded valence shell with one extra electron pair.
• LS =

Additional exercises on Lewis Structures are available in the Lewis Structure Module.

For a modern view of bonding illustrated with QuickTime movies based on quantum calculations you may enjoy the Supplement.

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